2. CHAPTER 13 – States of Matter

3. THE KINETIC THEORY 1.All matter is composed of very small particles 2.These particles are in constant, random motion

4. Kinetic Energy and Temperature “KINETIC” from a Greek word meaning “to move.” Kinetic Energy: Energy of an object in motion Temperature: A measure of the average kinetic energy of particles Increase Temperature, ________ Kinetic Energy Decrease Temperature, _______ Kinetic Energy Increase Decrease

5. Kinetic Energy of Gas Molecules m = mass v = velocity **If you increase the temperature, particles move FASTER.

6. The Four States of Matter This figure shows the four states of matter: _______, ________, ______ and __________(hydrogen nuclei and electrons). SolidLiquidGas Plasma

7. The Four States of Matter Solids: -Have a definite shape and volume -Particles vibrate about a fixed point Water is a solid below 0°C.

8. The Four States of Matter Liquids : -Flow, have a definite volume, and take the shape of its container -Particles are close together, but move randomly Water is a liquid between 0°C and 100°C.

9. The Four States of Matter Gases: -Takes the shape and volume of its container. -Are highly compressible. -Have very low densities. Water is a gas above 100°C. http://www.youtube.com/watch?v= s-KvoVzukHo

10. The Four States of Matter Plasma: -The fourth state of matter -Ionized gas or charged gas particles -Behaves differently from gases: (non-elastic collisions, attraction between particles, conducts electricity)

12. Kinetic Energy and Temperature If you continue to reduce temperature, what happens to the kinetic energy? All particle motion stops at o K or -273°C. Absolute zero (0 K, or –273°C) The temperature at which the motion of particles theoretically stops. Particles would have no kinetic energy (or motion). Absolute zero has NEVER been produced in the laboratory. http://www.youtube.com/watch?v=aOa14VQiu3Y https://www.youtube.com/watch?v=TNUDBdv3jWI

13. 1.Kinetic energy is the energy of ____? An object in motion 2.An increase in temperature increases/decreases the kinetic energy. increases 3.State of matter that has a fixed volume and takes the shape of its container. Liquid

14. 4. The temperature at which all motion stops? Absolute Zero 5.The state of matter made of ionized gas particles Plasma

15. Temperature Scales Fahrenheit: Water freezes at 32 ˚ F Water boils at 212 ˚ F Celsius (or Centigrade): Based on the freezing (0 ˚ C) and boiling points (100 ˚ C ) of water.

16. Temperature Scales Kelvin: 0 K is Absolute Zero No negative values in the Kelvin scale Temperature Conversions 0 K = -273 o C K  o C subtract 273 o C  K add 273 https://www.youtube.com/watch?v=TNUDBdv3jWI

17. CHECK: Convert…. 1.86 K to o C 2.58 o C to K 3.533 K to o C 4.-90 o C to K 86 K - 273 = -187 o C 58 o C + 273 = 331 K 533 K - 273 = 260 o C -90 o C + 273 = 183 K

18. What Causes Pressure of a Gas? Pressure = Force / Area Gas particles exert pressure when they collide with the walls of an object.

19. Pressure Atmospheric pressure is the weight of air per unit of area.

20. Measuring Pressure A barometer is the instrument we use to measure atmospheric pressure Atmosphere  1.00 atm = 760 mm Hg

21. Measuring Pressure A manometer is another instrument used to measure pressure

22. 1. kilopascals (kPa), 2. millimeters of mercury (mm Hg) 3. torr 4. pounds per square inch (lb/in 2 or psi) 5. atmosphere (atm) UNITS OF PRESSURE

23. Standard Pressure of a Gas Measured at sea level: 1 atm = 101.3 kPa = 760 mm Hg = 760 torr = 14.7 psi

24. Pressure Unit Conversions 1atm = 760 mmHg = 760 torr = 101.3kPa = 14.7psi Do the following conversion problems: 1.1.5 atm = kPa 2. 720 torr = atm 3. 75 kPa = mm Hg 4. 28.14 psi = atm

25. Standard Temperature and Pressure: STP the standard T & P for experimental measurements, to enable comparisons to be made between sets of data usually when working with gases. Standard temperature is 273 K or 0 °C Standard pressure is 1 atm or 101.3 kPa

26. Phase Changes Courtesy www.lab-initio.com

27. Energy Change of Phase Changes solidliquidgas Exothermic changes (kinetic energy or heat is released) Endothermic changes (kinetic energy or heat is absorbed) Deposition Sublimation Freezing Melting Condensing Boiling

28. Energy Change of Phase Changes solidliquidgas Exothermic changes (kinetic energy or heat is released or lost) Endothermic changes (kinetic energy or heat is absorbed or gained) Lose Gain Lose Gain Lose Gain

29. Chapter 13 Vocabulary so far…. Temperature Kinetic Energy Pressure Absolute Zero STP (Standard Temperature and Pressure)

30. MORE Chapter 13 Vocabulary Exothermic - kinetic energy or heat is released Endothermic - kinetic energy or heat is absorbed Normal Boiling Point – temperature at which a substance boils (or condenses) at Standard Pressure (1 atm or 760 mm Hg) Normal Freezing Point – temperature at which a substance freezes (or melts) at Standard Pressure (1 atm or 760 mm Hg)

31. Heat Curve of Phase Changes http://www.kentchemistry.com/links/Matter/HeatingCurve.htm

32. HEAT CURVE: Water phase changes Temperature remains __________ during a phase change. constant

33. Phase Diagram

34. MORE Chapter 13 Vocabulary Triple Point – T and P at which all three phases (solid, liquid, gas) are present Critical Point – T and P above which liquid and gas no longer exist as separate phases (above this, it is “supercritical fluid”)

35. Phase Diagram for Water

37. Carbon Phase Diagram for Carbon

38. Gases – Kinetic Molecular Theory 1.All matter is composed of very small particles 2. These particles are in constant, random motion 3. Collisions between gas particles are perfectly elastic 4. Gas particles are very small and are very far apart.

39. Kinetic Energy of Gas Molecules m = mass v = velocity **At any given temperature, the molecules of all gases have the same AVERAGE kinetic energy.

40. Kinetic Molecular Theory The average kinetic energy of the molecules is proportional to the absolute temperature.

41. Kinetic Energy of Gas Molecules Mixture of He (4 g/mol) and Ne (20 g/mol): Say the mixture has a KE = 2 J He: 2 J = 0.5 x (4 g) x (1m/s) 2 v = 1 m/s Ne:2 J = 0.5 x (20.2 g) x (0.44m/s) 2 v = 0.44 m/s Bigger gas particles move more __________. http://www.youtube.com/watch?v=UNn_trajMFohttp://www.youtube.com/watch?v=UNn_trajMFo (34 sec) Slowly

42. Kinetic-Molecular Theory Larger molecules (with more mass) move more slowly.

43. Diffusion of Gases Diffusion is the tendency of molecules to move toward areas of lower concentration until the concentration is uniform throughout.

44. Diffusion of Gases Diffusion is the tendency of molecules to move toward areas of lower concentration until the concentration is uniform throughout.

45. Kinetic-Molecular Theory Larger molecules (with more mass) move more slowly.

46. GRAHAM’S LAW OF DIFFUSION

47. The rate of effusion or diffusion is inversely proportional to the molar mass of the molecule. Mixture of He (4 g/mol) and Ne (20.2 g/mol): Which molecule moves faster?

48. GRAHAM’S LAW OF EFFUSION Example: What is the ratio of effusion rates for ammonia (NH 3 ) and hydrochloric acid (HCl)? NH 3 (17.0 g/mol) and HCl (36.5 g/mol): Which molecule moves faster?

49. GRAHAM’S LAW OF EFFUSION

50. Pressure Pressure definition: Force per unit area

51. Dalton’s Law of Partial Pressures In a mixture of gases, the total pressure is the sum of the partial pressures of the component gases.

52. Dalton’s Law – Partial Pressures The contribution each gas in a mixture makes to the total pressure is called the partial pressure exerted by that gas. P total = P A + P B + P C P total = 100 kPa + 250 kPa + 200 kPa P total = 550 kPa

53. Dalton’s Law – Another Problem A gas mixture containing oxygen, nitrogen, and carbon dioxide has a total pressure of 32.9 kPa. If P O2 = 6.6 kPa and P N2 = 23.0 kPa, what is P CO2 ? P total = P O2 + P N2 + P CO2 32.9 kPa = 6.6 kPa + 23.0 kPa + P CO2 P CO2 = 32.9 kPa – (6.6 kPa + 23.0 kPa) = 3.3 kPa

54. Real Gases In the real world, the behavior of gases only conforms to the ideal- gas equation at relatively high temperature and low pressure.

55. Deviations from Ideal Behavior The assumptions made in the kinetic-molecular model break down at high pressure and/or low temperature.

56. Deviations from Ideal Behavior NEED TO ADD – MORE DEVIATION FOR LARGE? MOLECULES AND POLAR MOLECULES

57. Avogadro’s Principle Use for gas stoichiometry problems – “Liter to Liter” conversions

58. Liter-Liter Conversions - Example What is the total number of liters of NO(g) produced when 1.6 liters of O 2 reacts completely with nitrogen at constant temperature and pressure? N 2 (g) + O 2 (g)  2 NO (g) Given (L of the known)# liters unknown = calculated L of unknown # liters known GIVEN MOLE RATIO FROM BALANCED CHEM EQN 1.6 L O 2 2 L NO = 3.2 L NO 1 L O 2

59. Liter-Liter Conversions - Example If I have 13 L of N 2 and excess H 2, how many liters of NH 3 can I make at constant temperature and pressure? N 2 (g) + 3 H 2 (g)  2 NH 3 (g) Given (L of the known)# liters unknown = calculated L of unknown # liters known GIVEN MOLE RATIO FROM BALANCED CHEM EQN 13 L N 2 2 L NH 3 = 26 L NH 3 1 L N 2